While the words “gas” and “vapor” are often used interchangeably, they represent two distinct physical states in chemistry and thermodynamics. Both describe a substance in a highly energetic, compressible phase that lacks a fixed volume or shape. The distinction is precisely defined by the relationship between a substance’s temperature and its unique thermodynamic properties.
Defining the States and Their Properties
A substance referred to as a gas is a phase that always expands to uniformly fill the entire volume of its container. Gases are characterized by weak intermolecular forces, allowing molecules to move rapidly and independently. A true gas cannot be forced into a liquid state simply by applying pressure, provided its temperature remains constant. This is because the kinetic energy of the molecules is too high for pressure to overcome the molecular motion and force them into a liquid phase.
A vapor describes a substance in the gas phase that is capable of coexisting with its liquid or solid form. Unlike a gas, a vapor can be condensed into a liquid by merely increasing the external pressure on it. This ability to easily revert to a liquid state is an observable difference between a vapor and a true gas. For example, water vapor is easily turned back into liquid water droplets through cooling or a slight pressure increase.
The Critical Point: The Thermodynamic Boundary
The scientific distinction between a gas and a vapor hinges on the critical point, a specific thermodynamic threshold. This point is defined by the critical temperature (\(T_c\)) and the critical pressure (\(P_c\)) for every pure substance. The critical temperature is the highest temperature at which a distinct liquid phase can exist. Above \(T_c\), the substance’s gaseous form cannot be liquefied, regardless of the pressure applied.
A substance in the gaseous state is classified as a vapor if its temperature is below \(T_c\). Because the temperature is low enough, an increase in pressure can overcome intermolecular forces, causing the substance to condense into a liquid. This allows a vapor to readily coexist in equilibrium with its liquid phase.
Conversely, a substance is defined as a gas if its temperature is above \(T_c\). When a gas is compressed above \(T_c\), it does not condense into a liquid. Instead, it transitions smoothly into a dense, fluid phase known as a supercritical fluid, where the boundaries between liquid and gas cease to exist.
Practical Applications and Common Usage
The critical point explains the common classification of substances. Water has a critical temperature of approximately \(374^\circ \text{C}\), far above standard room temperature. Since gaseous water in the atmosphere exists below this temperature, it is correctly referred to as water vapor. This is consistent with its ability to readily condense into liquid water upon cooling or compression.
In contrast, components of air like nitrogen and oxygen have much lower critical temperatures. Oxygen’s critical temperature is about \(-119^\circ \text{C}\), and nitrogen’s is around \(-147^\circ \text{C}\). Because ambient temperatures are always far above these values, these substances are perpetually above their critical temperatures and are classified as permanent gases.